1. Electronic Structure of Atoms

2. Waves
To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation.
The distance between corresponding points on adjacent waves is the wavelength ().

3. Waves
The number of waves passing a given point per unit of time is the frequency ().
For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.

4. Electromagnetic Radiation
All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.00  108 m/s.
Therefore,
c = 

5. The Nature of Energy
The wave nature of light does not explain how an object can glow when its temperature increases.
Max Planck explained it by assuming that energy comes in packets called quanta.

6. The Nature of Energy
Einstein used this assumption to explain the photoelectric effect.
He concluded that energy is proportional to frequency:
E = h
where h is Planck’s constant, 6.63  10−34 J-s.

7. The Nature of Energy
Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:
c = 
E = h

8. The Nature of Energy
Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.

9. The Nature of Energy
One does not observe a continuous spectrum, as one gets from a white light source.
Only a line spectrum of discrete wavelengths is observed.

10. The Nature of Energy
Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
Electrons in an atom can only occupy certain orbits (corresponding to certain energies).

11. The Nature of Energy
Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.

12. The Nature of Energy
Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by
E = h

13. The Wave Nature of Matter
Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties.
He demonstrated that the relationship between mass and wavelength was
 = h mv

14. The Uncertainty Principle
Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known:
In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
(x) (mv)  h 4

15. Quantum Mechanics
Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
It is known as quantum mechanics.

16. Energies of Orbitals
For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
That is, they are degenerate.

17. Pauli Exclusion Principle
No two electrons in the same atom can have exactly the same energy.
For example, no two electrons in the same atom can have identical sets of quantum numbers.

18. Electron Configurations
Distribution of all electrons in an atom
Consist of
Number denoting the energy level

19. Electron Configurations
Distribution of all electrons in an atom
Consist of
Number denoting the energy level
Letter denoting the type of orbital

20. Electron Configurations
Distribution of all electrons in an atom.
Consist of
Number denoting the energy level.
Letter denoting the type of orbital.
Superscript denoting the number of electrons in those orbitals.

21. Orbital Diagrams Each box represents one orbital.
Half-arrows represent the electrons.
The direction of the arrow represents the spin of the electron.

22. Hund’s Rule
“For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

23. Periodic Table We fill orbitals in increasing order of energy.
Different blocks on the periodic table, then correspond to different types of orbitals.

24. Some Anomalies
Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

25. Some Anomalies For instance, the electron configuration for copper is
[Ar] 4s1 3d5
rather than the expected
[Ar] 4s2 3d4.

26. Some Anomalies
This occurs because the 4s and 3d orbitals are very close in energy.
These anomalies occur in f-block atoms, as well.

27. Paramagnetic and Diamagnetic
Subshells are not completely filled
Example:
He:
Be:
Li: N:
Li and N are Paramagnetic
He and Be are Diamagnetic

28. Isoelectric Having the same electron configuration Examples:
Lithium ion and Helium Atom
Strontium ion and Rubidium Ion and Krypton